Oxidation-reduction potentials for plutonium. Oxidation reduction potentials [Al] for plutonium ions are given in Table 9.5..

Other potentials may be obtained as linear combinations of these. For instance,

Pu4+ + 2H20 ->■ Pu022+ + 4ІГ + 2e’

Ett = — y’— = -1.0433 V (9.29)

Equilibria for the reduction of Pu02* or Pu02 2+ to Pu3+ or Pu4+ depend on the fourth power of the hydrogen ion concentration. Increasing acidity displaces equilibrium toward the reduced states.

If the salts and acids in a given system were completely dissociated into simple ions with no complex formation and if the activity coefficients were always the same, oxidation-reduction potentials would be independent of the anionic or cationic species present. However, as we saw in Sec. 1.1, these conditions are rarely obtained with the higher-valent actinides. Therefore, oxidation-reduction potentials for the actinides are usually measured in perchlorate solutions, where relatively little complexing and a high degree of dissociation occur [Al, SI]. These are referred to as “formal potentials,” which are more accurate for practical calculations. To obtain standard potentials these formal potentials, or equivalent thermodynamic data, must be extrapolated to zero ionic strength. Formal potentials for other solutions, such as 1 M HN03 or 1M NaOH, are also reported in the literature [А1]. As pointed out in Sec. 1.3, in a practical system such as nitric acid solutions the equilibria estimated from oxidation-reduction potentials must be corrected for complexing, as well as incomplete dissociation. For example, the oxidation potential for the couple Pu(III)-Pu(IV) is—0.9819 V in 1 M HC104 but —0.92 V in 1 M HNOj

[SI]-

Oxidation-reduction potentials of the actinides. The formal potentials for transition between the valence states of the actinides are listed in Table 9.6.

The stability of an intermediate oxidation number against disproportionation can be obtained as follows. Consider the disproportionation of U(V) according to the following reaction:

Uv02+ -*■ UV1C>22+ + e" £l,6 =-0.063 V (930)

Ц^ + гНіО-* Uv02* + 4H+ + e’ £4,5 = -0.613 V (931)

2Uv02+ + 4H+-UVI022t + U4+ + 2H20 AE = -0.063 -(-0.613) = 0.550 V (9.32) The equilibrium constant for the overall reaction (9.32) is

£ = e38.93(-0.063 + 0.613) = 1 99 X 109 (!) .33)

Hence, pentavalent uranium is unstable in aqueous solution at [H+] > 1.

Disproportionation of valence n to valences n + 1 and n — 1 will proceed spontaneously at pH = 0 if the potential for the oxidation from n to n + 1 is larger or more nearly positive than the potential for oxidation from n — 1 to n. Applying this criterion to the data of Table 9.6,

Table 9.5 Oxidation-reduction potentials for plutonium

we see that in the tetravalent state uranium, neptunium, and plutonium are stable. In the pentavalent state protactinium, neptunium, and americium are stable (cf. Table 9.4).

The positive potential for U(III)-U(IV) indicates that the unstable U(III) would be rapidly oxidized by water in aqueous solution. The relatively low negative potentials for the oxidation of U(III) through intermediate states to U(VI) indicate that the latter should be quite stable in aqueous solutions. The elements of higher atomic number become progressively more difficult to oxidize to the hexavalent state.

Solutions containing U(V1) and Pu(III) or Pu(IV), as used in aqueous separation processes, are stable against oxidation of plutonium by uranium because the potentials for the transitions U(TV) to U(VT) and U(V) to U(VI) are more nearly positive than the plutonium potentials. Plutonium may be reduced from Pu(IV) to Pu(III) without affecting uranium oxidation by choosing a reducing agent, such as Fe2+, whose oxidation potential is less negative than the —0.9819 V required for Pu(IV) reduction and more negative than the —0.338 V that would reduce U(VI).

The data indicate that U(VI) should oxidize Np(III) to Np(IV).

Oxidation-reduction potentials for couples consisting of the actinides or the fission products in acid solution (1 M HC104) are listed in Table 9.7. Potentials for a selected group of oxidizing and reducing agents are listed in Table 9.8. The couples are listed in order of decreasing strength as reducing agents. In the cases where the molecular and ionic species involved in a given valence transition are different in acidic and basic solutions, the acid system (1 M НСЮ4) has been chosen.

The oxidation-reduction schemes of the more important multivalent elements encountered in aqueous fuel reprocessing are summarized in Fig. 9.1.

Rate of oxidation-reduction reactions. Oxidation-reduction reactions that involve only the transfer of electrons from one uncomplexed ion to another in an ionizing solvent are reversible and, for all practical purposes, instantaneous. Equation (9.8) is an example. On the other hand, reactions involving molecular rearrangements, even though thermodynamically possible, may be

very slow or may not go at all. A familiar example of a slow reaction is the gradual approach to the end point in titration of ferrous ion with permanganate ion in acid solution:

8H4 + 5Fe2+ + Mn04‘ -*■ 5Fe34 + Mn24 + 4H2 О (9.34)

The oxidation of an actinide ion from M3+ or M4+ to M024 or M0224, or its reduction from M(V or VI) to M(IV or III), is inconveniently slow, apparently because of the sluggish combination of M and О in the oxidation step or the slow breaking of the M—О bond in the reduction step. Of the three couples involving plutonium, Eq. (9.27) is very slow, whereas (9.26) and (9.28) are practically instantaneous. Further discussion of the rate of oxidation — reduction of plutonium solution appears in Sec. 4.6.